Module 7 Lesson 3 - Read
Read: Size of Atoms and Ions
Overview
Learning the ins and outs of the periodic table is essential for a deeper understanding of chemistry. Effective nuclear charge, shells, and shields affect all trends on the periodic table. After understanding those concepts, the periodic trends can be learned. This lesson will look specifically at atomic size and ionic size.
Finding the Periodic Trends
Effective nuclear charge, Zeff, is the attraction of protons within the nucleus to its outer valence electrons within many-electron atoms. If there are multiple electrons in an atom, there is not only an attraction of electrons to the nucleus but also electron-electron repulsion. Some of the electron-electron repulsion cancels out the electron-nucleus attraction. Effective nuclear charge can be used to solve for the actual pull an electron feels from the nucleus in a many-electron atom. Effective nuclear charge helps explain all trends on the periodic table that move across a period (left to right or right to left).
Shell is a synonym for energy level. As the number of energy levels increases, the number of shells increases, meaning there are more core (inner) electrons shielding the outer valence electrons. This makes it more difficult for an outer valence electron to feel attraction from the nucleus. The shielding effect helps explain all trends on the periodic table that move up or down a column.
Atomic Size
There are two ways atoms can come in contact with each other. They can collide and not bond or they can collide and bond. Different ways of explaining about the radii for these two situations exist.
A nonbonding atomic radius occurs when two atoms collide but do not bond. This is also called Van der Waals radius. The two atoms squeeze together as closely as they can until their electron clouds cannot overlap due to electron-electron repulsion between the two atoms. A nonbonding radius is the shortest length separating the nuclei of the two atoms during a collision. This is twice the radii of the atoms involved in the collision.
To find the nonbonding radius, add the atomic radius 1 and atomic radius 2 together.
The other situation occurs when two atoms collide and then bond. Two atoms adjacent to each other can have an attractive interaction, and this is referred to as a chemical bond. When this happens, these two atoms are closer together than in a nonbonding collision. Half the distance between two bonding atoms’ nuclei is called the bonding atomic radius.
The following picture shows the distance between the two nuclei, d. The bonding atomic radius is half the distance of d. Thus, the bonding atomic radius is
d.
Scientists have discovered different ways of measuring bonding atomic radius for all elements on the periodic table. They observe the distances of these atoms when they bond in common molecules. A chart of average atomic radii was made for each element as a way to help predict what length each would be when bonded with another atom.
The trend for bonding atomic radii decreases from left to right across a period and increases from top to bottom in a column on the periodic table. The elements with the smallest bonding atomic radii are located in the top right of the periodic table. The elements with the largest bonding atomic radii are located in the bottom left of the periodic table.
Ionic Size
The general trend for ionic size is similar to the general trend for bonding atomic radii.
The first step is to review what an ion is and the two types of ions that can be made. An ion occurs when a neutral atom has gained or lost electrons, creating an electric charge. This charge creates attraction and/or repulsion between ions. An atom that has lost electrons creating a positive charge is called a cation. An atom that has gained electrons creating a negative charge is called an anion.
Cations are smaller than their neutral atoms. When an electron is lost, the effective nuclear charge becomes larger and has the strength to pull the remaining electrons closer to the nucleus, thus making the cation’s size smaller than before.
Anions are larger than their neutral atoms. When an electron is gained, the effective nuclear charge becomes smaller. The nucleus has more electrons to attract than it did before, making it more difficult. The resulting anion is now larger than its neutral counterpart.
The ionic radius size decreases from left to right across a period and increases from top to bottom in a column on the periodic table.
Expand: Which is Larger? Which is Smaller?
Overview
Looking at the periodic table and knowing how to predict trends is the next step in concept understanding. This section asks learners to predict size based on location on the periodic table for neutral atoms and the ionic size.
Predicting Bonding Atomic Radii Size
Example 1:
Place the following atoms in order of increasing (smallest to largest) atomic size: Cl, F, S, and O.
Answer: Smallest F < O < Cl < S Largest. According to the periodic trend, atomic size increases from right to left and top to bottom. Fluorine has the smallest atomic size due to having the largest effective nuclear charge. Its electrons are pulled closer to the nucleus making its atomic size very small. Oxygen is the next largest because it has one less proton in its nucleus. Chlorine is the second largest because it has one more shell (energy level) than fluorine and oxygen. Sulfur has the largest atomic size due to its extra shell and weaker effective nuclear charge.
Example 2:
Which of the following has the largest atomic size? Ca, Na, K, Mg
Answer: The largest is potassium (K). Potassium has one more energy level than sodium and magnesium. Potassium has a weaker effective nuclear charge than calcium.
Example 3:
Which of the following has the smallest atomic size? Ca, Na, K, Mg
Answer: The smallest is magnesium. Magnesium has the smallest number of energy levels and the strongest effective nuclear charge.
Predicting Ionic Radii Size
Example 4:
Which is larger: Cl or Cl-1?
Answer: Cl-1 is larger. The anion is larger than its neutral atom. Adding an extra electron makes the nucleus less effective at attracting electrons. Since the nucleus is unable to “hold on to” its electrons as effectively, the electrons exist at a further distance than they did before. This makes the atomic size larger.
Example 5:
Which is larger: K or K+1?
Answer: K is larger. When the potassium atom loses an electron, the nucleus has fewer electrons to attract than it did before. This allows the nucleus to “pull in” those remaining electrons more closely making the atomic size smaller than it was originally.
Module 7 Lesson 3 of 5